Thermochemistry!

Thermochemistry! AP Chapter 5

Temperature vs. Heat • Temperature is the average kinetic energy of the particles in a substance. • Heat is the energy that is transferred from one object to another. • Heat always flows from the hotter object to the colder object.

Energy! • Energy is the ability to do work. • Kinetic Energy- the energy of motion • Potential Energy – the energy that an object has as a result of its composition or its position with respect to another object.

Units of Energy • 1 Joule = 1 kg m2/s2 (1 kJ is 1000J) • Used to calculate the energies associated with chemical reactions. • Calorie – Amount of energy required to raise the temperature of 1 gram of substance 1 °C. (This is specific heat!) 1 calorie will raise the temperature of 1 g of H2O from 14.5 °C to 15.5 °C. 1 calorie is equal to 4.184 Joules (exactly!)

Systems and Surroundings • System – the portion used in a study. • It can be an open system or a closed system. • Opensystem – matter and energy can interact with the surroundings. • Closedsystem – the matter cannot interact with the surroundings.

First Law of Thermodynamics • Energy Is Conserved!

Internal Energy • Internal Energy is the sum of all the kinetic and potential energies of all its components. ΔE = Efinal - Einitial

ΔE • A positive value for ΔE is when Efinal > Einitial • If energy has been absorbed from its surroundings, it is endothermic. • If energy is given off to the surroundings, it is exothermic.

Initial state refers to the reactants, while final state refers to the products.

Endothermic reaction Exothermic reaction

A system composed of H2 (g) and O2 (g) has greater internal energy than a system composed of H2O (l). Gases have greater kinetic energy and must lose some of that energy to change states back to the liquid state.

Internal energy is a function of state.

If a battery is shorted out and loses energy to the environment only as heat, no work is done. • If a battery is discharged and loses energy as work (to make the fan run) it also loses heat energy. • The value of ∆E is the same.

Enthalpy • The change in enthalpy for a reaction (∆H) is the overall measure of energy that is absorbed to break bonds and the energy that is released when new bonds form. • A reaction is said to be spontaneous if it occurs without being driven by an outside force. (driving forces are enthalpy & entropy) • ∆H = ΣH(products) - ΣH(reactants)

In an endothermic system where it absorbs heat, ∆H will be positive (∆H > 0). In an exothermic system, where heat is given off, ∆H will be negative (∆H < 0).

Enthalpy Diagrams • Enthalpy is an extensive property – it depends on how much you have. If 1mol of CH4 and 2 mol O2 yield -890 kJ, then 2 mol CH4 and 4 mol O2 would yield double that. • The enthalpy change for a reaction is equal in magnitude, but opposite sign, for a reverse reaction.

Calorimetry • This is a measure of the amount of energy that is needed or lost when a certain mass of a substance changes temperature. • q = mC∆T • q is the amount of energy (J) • m is the mass of the substance (g) • C is the specific heat capacity of the substance • ∆T is the change in temperature

Calorimeters • Calorimeters are devices that measure the transfer of heat from one object to another.

Heat of Formation (∆H°f) • The heat change that occurs when one mole of a compound is formed from its elements at 1 atm pressure. • Generally, the standard enthalpy of formation for any element in its most stable form is 0. (i.e. O2 gas would have a standard enthalpy of 0.) • Remember Appendix C!

Standard Enthalpy Changes • The standard enthalpy change can be calculated from the standard enthalpies of formation of the reactants and products in the reaction (see Appendix C for values.) • ∆H°rxn = Σn∆H°f(products) - Σm ∆H°f(reactants) • The n and m refer to the molar coefficients in the chemical equation.

Also refer to Appendix C!

Hess’s Law • If you can break a chemical reaction into several steps, add up all of the ∆H’s for each step to get the overall ∆H for the reaction.

Entropy • Entropy is a measure of randomness or disorder of a system. The greater the disorder, the greater the entropy. • In terms of entropy, gases>liquids>solids. • When pure substance dissolves in a liquid, its entropy increases. • When gas molecules escape a solvent, entropy increases. • Entropy increases with molecular complexity. • Reactions that increase the number of moles of particles often increase the entropy of the system.

Predict! • Na+(aq) + Cl-(aq) → NaCl (s) ∆S is negative • NH4Cl (s)→ NH3(g) + HCl (g) ∆S is positive

Thermochemistry!